• An atom is the basic building
block of matter.
• All objects are made of
atoms. The air, your desk and
all living things are made of
atoms.
• Atoms are EXTREMELY small.
One atom is only one ten-
billionth of a meter wide!
© Stephanie Elkowitz 1 Atoms & Reactions
What is an
atom?
• The idea of an atom was
first developed around 460
B.C. by Green philosopher
Democritus.
• Democritus believed that
you could not indefinitely
break an object in half. At
some point you get to the
smallest bit of matter, which
cannot be broken. He called
this bit of matter an atom.
© Stephanie Elkowitz 2 Atoms & Reactions
• In 1805, English chemist John
Dalton proposed and
published he theory on atoms.
• Dalton believed matter was
made of extremely small
atoms. Atoms of the same
substance have identical size,
mass and other properties.
• Dalton also believed atoms
could not be created,
subdivided or destroyed but
could combine to form
different chemical substances. © Stephanie Elkowitz 3 Atoms & Reactions
• In 1897, English physicist J.J.
Thomson discovered the
electron.
• Thomson discovered that an
electron is a tiny, negatively
charged particle.
• From this finding, Thomson
proposed the first model of an
atom.
© Stephanie Elkowitz 4 Atoms & Reactions
• J.J. Thomson came up with his
model of the atom in 1904.
• According to Thomson, an atom
was made of negatively charged
particles (electrons) embedded in a
“soup” of positive charges.
• This model suggested atoms
resembled plum pudding. The
electrons are “plums” surrounded
in a positively charged “pudding.”
For this reason, Thomson’s model
is called the plum pudding model.
© Stephanie Elkowitz 5 Atoms & Reactions
• Ernest Rutherford proposed an
alternative model in 1911.
• According to Rutherford, an atom is
made of a central, positively
charged region. He called this
region the nucleus.
• Rutherford believed electrons
surrounded the nucleus as a
“cloud.”
• He also believed the nucleus of the
atom was small and dense
compared to the overall size of the
atom. © Stephanie Elkowitz 6 Atoms & Reactions
• Rutherford proved his model of the
atom with a famous experiment
known as the Gold Foil
Experiment.
• Rutherford worked with two other
scientists on this experiment: Han
Geiger and Ernest Marsden.
• In this experiment, Rutherford shot
a bean of positively charged
particles, called alpha particles, at
a thin piece of gold foil. He
recorded where the alpha particles
scattered as they struck the gold. © Stephanie Elkowitz 7 Atoms & Reactions
Rutherford, Geiger & Marsden
• If Thomson’s model was correct,
all the alpha particles should pass
straight through the foil as the
particles struck the gold.
• So, why would they pass straight
through? According to Thomson’s
plum pudding model, the charges
are symmetrically and evenly
distributed through an atom. This
is why the alpha particles would
pass straight through.
© Stephanie Elkowitz 8 Atoms & Reactions
• Rutherford observed most particles
pass straight through. However, some
particles scattered and some particles
“bounced off” the gold foil.
• Rutherford believe his atomic model
explained these experimental results:
– Most particles passed through because an
atom is mostly empty space.
– Some particles scattered because they
were deflected by negatively charged
electrons.
– Some particles reflected back because
they bounced off the positively charged
nucleus. © Stephanie Elkowitz 9 Atoms & Reactions
• In 1914, Neils Bohr modified
Rutherford’s model of the atom.
• Like Rutherford, Bohr believed
an atom is made of a small,
positively charged nucleus
surrounded by negatively
charged electrons.
• However, Bohr believed the
electrons traveled around the
nucleus in circular orbits.
• This model is known as the
Rutherford-Bohr Atomic Model.
© Stephanie Elkowitz 10 Atoms & Reactions
• Today we refer to the
Rutherford-Bohr Atomic
Model when studying atoms.
• According to this model,
there are two major regions
of an atom:
1. The center region of the
atom is called the nucleus.
2. The surrounding area around
the nucleus is mostly empty
space. Electrons orbit the
nucleus within this area.
© Stephanie Elkowitz 11 Atoms & Reactions
Nucleus
• Atoms are VERY tiny! They are so small that we use special
units of measurement to describe the size and mass of
atoms and particles within atoms.
• We measure the mass of atoms in atomic mass units (u).
• One atomic mass unit equals 1.66 x 10-27 kilograms.
• We measure the size of atoms in picometers (pm).
• One picometer equals one trillionth of a meter. In other
words, 1 picometer equals 1.00 x 10-12 meters.
• An atom is approximately 100 pm, or 1 ten-billionth of a
meter wide.
© Stephanie Elkowitz 12 Atoms & Reactions
• Subatomic particles are the
particles found within an
atom.
• There are three subatomic
particles:
– Proton (p)
– Neutron (n)
– Electron (e-)
We often use abbreviations to
denote the different subatomic
particles. The abbreviations are
listed next to the particles above.
© Stephanie Elkowitz 13 Atoms & Reactions
Proton Neutron
Electron
• Protons are found in the
nucleus of an atom.
• A proton is positively
charged.
• One proton has a +1
charge.
• A proton has a mass of 1
atomic mass unit.
© Stephanie Elkowitz 14 Atoms & Reactions
Proton +
+
• Neutrons are also found in
the nucleus of an atom.
• A neutron is a neutral
particle. On other words, a
neutron has NO charge.
• A neutron has a mass of 1
atomic mass unit
© Stephanie Elkowitz 15 Atoms & Reactions
Neutron
• Electrons are found in the
space surrounding the
nucleus.
• An electron is negatively
charged. One electron has a
-1 charge.
• An electron has negligible
mass. In other words, the
mass of an electron is SO
small, it is insignificant. To
keep it simple, we say
electrons have no mass.
© Stephanie Elkowitz 16 Atoms & Reactions
Electron
–
–
• Because the nucleus
contains positive and
neutral particles, the net
charge of the nucleus is
POSITIVE.
• A force known as the
nuclear force holds the
proton(s) and neutron(s)
together in the nucleus.
© Stephanie Elkowitz 17 Atoms & Reactions
• The space around the
nucleus is NEGATIVELY
charged because it contains
negatively charged
electrons.
• Electrons stay in orbit
around the nucleus due to
electromagnetic force. This
force “holds” an atom
together.
© Stephanie Elkowitz 18 Atoms & Reactions
• Although we cannot view the
arrangement of particles
within an atom, we have an
instrument that allows us to
see the surface of atoms.
• A scanning tunneling
microscope (STM) can view
the atoms of the surface of an
object.
• This image was taken with an
STM. It shows the surface of a
piece of gold. You can actually
see individual atoms!
© Stephanie Elkowitz 19 Atoms & Reactions
• All atoms have the same arrangement of subatomic
particles. A change in the amount of subatomic particles
will change the type of atom.
• An element is a type of atom with the same number of
protons.
• We also use the term element to describe a substance that
is made of the same type of atom.
• There are more than 100 different types of elements.
• An element is abbreviated with a one or two letter symbol.
© Stephanie Elkowitz 20 Atoms & Reactions
• Examples of elements and their symbols are listed below:
© Stephanie Elkowitz 21 Atoms & Reactions
Element Symbol
Hydrogen H
Carbon C
Oxygen O
Nitrogen N
Sodium Na
Lithium Li
Element Symbol
Neon Ne
Helium He
Fluorine F
Chlorine Cl
Aluminum Al
Iron Fe
• Electrons constantly move
around the nucleus of an
atom.
• They orbit the nucleus in
specific orbitals or shells that
surround the nucleus.
© Stephanie Elkowitz 22 Atoms & Reactions
+
• A shell is also called an energy level.
• A shell is called an energy level
because it is associated with a
certain amount of energy.
• The shell closest to the nucleus has
the lowest energy. Electrons in the
first shell have the least amount of
energy.
• As you move away from the nucleus,
the energy associated with a shell
increases. Electrons in the shell
furthest away from the nucleus have
the most energy. © Stephanie Elkowitz 23 Atoms & Reactions
INCREASING
ENERGY
• Each shell has a
maximum amount of
electrons it can hold.
– The 1st shell holds 2 e-.
– The 2nd shell holds 8 e-.
– The 3rd shell holds 18 e-.
– The 4th shell holds 32 e-.
• In large atoms, you can
find up to 7 shells.
• No shell can hold more
than 32 electrons.
© Stephanie Elkowitz 24 Atoms & Reactions
+
2 e-
8 e-
18 e-
32 e-
• How electrons are
arranged in the shells of
an atom is called
electron configuration.
• Electrons “fill” or take up
space in each shell.
• In general, electrons fill
lower shells first
because lower shells
have less energy.
• Once a shell is full,
electrons begin to fill the
next higher shell. © Stephanie Elkowitz 25 Atoms & Reactions
Fill this shell first...
then this shell...
and finally this shell.
• To show electron configuration, we
draw a Bohr Diagram.
• To draw a Bohr Diagram:
1. Draw a circle to represent the
nucleus of the atom.
2. Write the element’s symbol, number
of protons (p) and number of
neutrons (n) inside the circle.
3. Draw rings around the circle to
represent electron shells. Each ring
represents a different energy level.
4. Draw electrons as dots in the rings.
Remember, each “ring” can only hold
so many electrons.
© Stephanie Elkowitz 26 Atoms & Reactions
Example: Oxygen (O)
8 protons, 8 neutrons, 8 electrons
Example 1: Draw a Bohr diagram of Carbon (C) that has 6
protons, 6 neutrons and 6 electrons.
© Stephanie Elkowitz 27 Atoms & Reactions
Remember...
2 e- fill the 1st
shell, 8 e- fill
the 2nd shell
and 18 e- fill
the 3rd shell.
Example 1: Draw a Bohr diagram of Carbon (C) that has 6 protons,
6 neutrons and 6 electrons.
© Stephanie Elkowitz 28 Atoms & Reactions
2 electrons completely fill the 1st shell
4 electrons partially fill the 2nd shell
C
6 p
6 n
Example 2: Draw a Bohr diagram of Sodium (Na) that has 11 protons,
12 neutrons and 11 electrons.
© Stephanie Elkowitz 29 Atoms & Reactions
Example 2: Draw a Bohr diagram of Sodium (Na) that has 11 protons,
12 neutrons and 11 electrons.
© Stephanie Elkowitz 30 Atoms & Reactions
Na
11 p
12 n
• The electrons found in the
outermost orbital are called
valence electrons.
• For this reason, the
outermost shell is called the
valence shell.
• The number of valence
electrons determines many
chemical properties of an
element.
© Stephanie Elkowitz 31 Atoms & Reactions
valence
electrons
• An atom cannot have more
than 8 valence electrons.
• An atom with 8 valence
electrons is said to have a
full outer shell.
• For example, neon (Ne) has
8 valence electrons.
© Stephanie Elkowitz 32 Atoms & Reactions
• Helium (He) has a full
valence shell with only two
electrons.
• Helium only has one shell.
The maximum amount of
electrons held in the first
shell is 2 electrons. Since
the shell holds the
maximum amount of
electrons it can hold, helium
is said to have a full valence
shell.
© Stephanie Elkowitz 33 Atoms & Reactions
• There are two important numbers associated with an
element that help you determine the number of protons,
neutrons and electrons in a neutral atom of that element:
1. Atomic Number
2. Atomic Mass
© Stephanie Elkowitz 34 Atoms & Reactions
• The atomic number is the number of protons found in an
element.
• Atoms of the same element have the same number of
protons and thus, the same atomic number.
• Example: The atomic number of Helium is 2. Therefore, there
are two protons in an atom of Helium.
© Stephanie Elkowitz 35 Atoms & Reactions
• Atomic mass is the mass of an atom. It is measured in atomic
mass units (u).
• Atomic mass equals the sum of protons and neutrons in an
atom.
• Remember...
– The mass of one proton is 1 u.
– The mass of one neutron is 1 u.
– The mass of an electron is negligible. In other words, an
electron has insignificant mass and it does not contribute
to the weight of an atom.
© Stephanie Elkowitz 36 Atoms & Reactions
Example: Carbon has 6 protons and 6 neutrons. What is
carbon’s atomic mass?
Protons + Neutrons = Atomic Mass
6 + 6 = 12
Atomic Mass = 12 atomic mass units
© Stephanie Elkowitz 37 Atoms & Reactions
• If you know an atom’s atomic number and the number of its
neutrons, you can calculate atomic mass.
• Adding an element’s atomic number (which is equivalent to
the number of protons) to the number of neutrons equals the
elements atomic mass.
© Stephanie Elkowitz 38 Atoms & Reactions
• If you know an atom’s atomic mass and atomic number, you
can calculate the number of neutrons in an atom.
• To find the number of neutrons, subtract atomic number from
the atomic mass.
© Stephanie Elkowitz 39 Atoms & Reactions
Example: Nitrogen has an atomic mass of 14. Its atomic number
is 7. How many neutrons are found in this atom?
Atomic Mass − Atomic Number = # of neutrons
14 − 7 = 7
Number of neutrons = 7
© Stephanie Elkowitz 40 Atoms & Reactions
• The number of protons is the same for all atoms of a
specific element.
• A change in the number of protons will change the type of
atom - the element to which the atom belongs.
• The number of neutrons and electrons in an atom can vary
without changing the element to which the atom belongs.
• A change in the number of neutrons will change the form of
an atom. Different forms of atoms are called isotopes.
• The change in the number of electrons will change the
electric charge of an atom. Atoms with an electric charge
are called ions.
© Stephanie Elkowitz 41 Atoms & Reactions
• An isotope is a variety of an element with a different number
of neutrons.
• The name of an isotope is the name of the element followed
by a dash (-) and the atomic mass of the isotope.
• Example: Carbon can have 6, 7 or 8 neutrons. Carbon with
6 neutrons is called Carbon-12. Carbon with 7 neutrons is
called Carbon-13. Carbon with 8 neutrons is called Carbon-
14.
© Stephanie Elkowitz 42 Atoms & Reactions
• The average atomic mass is the average of all the naturally
occurring isotopes of an element.
• To find average atomic mass:
1. Identify the different isotopes of the element and the
atomic mass of each isotope.
2. Multiple the atomic mass of each isotope by its percent
abundance (in decimal form). Percent abundance is the
percent the element is found in the natural world.
3. Find the sum of these values. The sum is equal to the
average atomic mass.
© Stephanie Elkowitz 43 Atoms & Reactions
Example: Chlorine
Step 1: Identify the different
isotopes of the element and the
atomic mass of each element.
Step 2: Multiple the atomic mass of
each isotope by its percent
abundance (in decimal form).
Step 3: Add these values together
to find average atomic mass.
© Stephanie Elkowitz 44 Atoms & Reactions
Chlorine-35 exists 76.77% abundance
Chlorine-37 exists 24.23% abundance
.7577 x 35 = 26.5195
.2423 x 37 = 8.9651
26.5195 + 8.9651 = 35.4846 ≈ 35.48
• Average atomic mass gives you an idea as to what the most
common isotope of an element is.
• To find the most common isotope from the average atomic
mass, round the average atomic mass to the nearest whole
number. This “trick” works for most elements.
• For example, the average atomic mass of carbon is 12.011.
Because this value is closest to 12, we can assume that the
most common isotope of Carbon is Carbon-12.
© Stephanie Elkowitz 45 Atoms & Reactions
• Most isotopes are stable. Stable isotopes have a “happy”
balance of protons and neutrons.
• Some isotopes are not stable. An unstable isotope is called a
radioactive isotope. A radioactive isotope is called so
because it emits radiation.
• Radiation is the release of energy in the form of waves or
subatomic particles. Radiation is dangerous because it can
harm the cells of living things. Specifically, it can alter
genetic material (DNA) in cells.
© Stephanie Elkowitz 46 Atoms & Reactions
• A radioactive isotope releases energy in order to become
stable. This process is called radioactive decay.
• During radioactive decay, the atom emits radiation.
Radiation can be high-energy waves and/or subatomic
particles.
• As the atom undergoes radioactive decay, it becomes stable.
As a stable isotope, the atom no longer emits radiation.
© Stephanie Elkowitz 47 Atoms & Reactions
• An element can gain or lose electrons to form an ion.
• An ion is an atom with an electric charge.
• Neutral atoms have the same amount of protons and
electrons. These atoms have zero net charge.
• An ion does not have the same amount of protons and
electrons. These atoms have a positive or negative change.
© Stephanie Elkowitz 48 Atoms & Reactions
• When an atom loses an electron, it loses a negative charge.
Therefore, the atom has more protons and is positive.
• A positive ion is called a cation.
© Stephanie Elkowitz 49 Atoms & Reactions
• When an atom gains an electron, it gains a negative charge.
Therefore, the atom has more electrons and is negative.
• A negative ion is called an anion.
© Stephanie Elkowitz 50 Atoms & Reactions
• To determine the electric charge of an ion, you must know
the difference between protons and electrons.
• If there are more protons, the ion is positive A “+” is used to
denote a positive ion.
• If there are more electrons, the ion is negative a “−” is used
to denote a negative ion.
• Write the charge as a superscript to the right of the atom.
– If the ion has a +1 charge, simply write +.
– If the ion has a −1 charge, simply write −.
© Stephanie Elkowitz 51 Atoms & Reactions
Example: A sodium atom has 11 protons and 10 electrons. How
do you denote the charge on this atom?
© Stephanie Elkowitz 52 Atoms & Reactions
Example: A sodium atom has 11 protons and 10 electrons. How
do you denote the charge on this atom?
1. 11 − 10 = 1 (the different between protons and electrons)
2. There are more protons, so the ion is positive (+).
3. Na+ is the abbreviation for a sodium ion with a +1 charge.
© Stephanie Elkowitz 53 Atoms & Reactions
Example: An oxygen atom has 8 protons and 10 electrons. How
do you denote the charge on this atom?
© Stephanie Elkowitz 54 Atoms & Reactions
Example: An oxygen atom has 8 protons and 10 electrons. How
do you denote the charge on this atom?
1. 10 − 8 = 2 (the different between protons and electrons)
2. There are more electrons, so the ion is negative (−).
3. O−2 is the abbreviation for an oxygen ion with a −2 charge.
© Stephanie Elkowitz 55 Atoms & Reactions
• An atom wants to gain or
lose electrons in order to
have a complete valence
shell.
• Remember... A complete
valence electron shell holds
8 electrons except for the
first shell - the first shell can
hold a maximum of 2
electrons
© Stephanie Elkowitz 56 Atoms & Reactions
Why do
atoms
become
ions?
• An atom that has only one or two valence electrons tends to
lose those electrons and become a positive ion.
– Example: Lithium (Li) is an atom with 1 valence electron. Lithium
tends to lose this electron to become a +1 ion with a complete
valence electron shell.
© Stephanie Elkowitz 57 Atoms & Reactions
• An atom that has nearly a full valence shell tends to gain
electrons and become a negative ion.
– Example: Fluorine (F) has 7 valence electrons. It tends to gain one
electron to become a −1 ion with a complete valence electron shell.
© Stephanie Elkowitz 58 Atoms & Reactions
• There are more than 100 different elements. All the known
elements are organized in a table. This table is called the
Periodic Table of Elements.
© Stephanie Elkowitz 59 Atoms & Reactions
• The periodic table displays
each element in a box.
• In most period tables, each
box gives 4 pieces of
information about an element:
①The name of the element
② The chemical symbol of the
element
③ The atomic number
④ The average atomic mass
© Stephanie Elkowitz 60 Atoms & Reactions
1
H Hydrogen
1.00794
①
④
②
③
• Elements are presented in the periodic table by order of
increasing atomic number.
© Stephanie Elkowitz 61 Atoms & Reactions
• The elements are arranged in rows and columns.
• The rows are called periods. The columns are called groups.
© Stephanie Elkowitz 62 Atoms & Reactions
• Electron configuration
dictates how elements are
organized in the rows and
columns of the period table.
In general:
– Elements with the same
number of valence electrons
are placed in the same column.
– Elements with the same
number of shells are placed in
the same row.
• This organization explains the
shape of the periodic table.
© Stephanie Elkowitz 63 Atoms & Reactions
Why is the
periodic table
shaped the
way it is?
• There are four major types of elements in the periodic table:
1. Metals 3. Metalloids
2. Nonmetals 4. Noble Gases
© Stephanie Elkowitz 64 Atoms & Reactions
• Most elements are metals.
• Metals are hard, shiny (lustrous), malleable and ductile.
• Metal are good conductors and usually have a high density
and melting point.
• Examples: Sodium, Calcium, Iron, Gold
© Stephanie Elkowitz 65 Atoms & Reactions
• Nonmetals make up the majority of the universe.
• Nonmetals are brittle and non-elastic. Most are dull.
• Nonmetals are poor conductors and usually have a low
melting point.
• Examples: Carbon, Phosphorus, Sulfur, Iodine, Chlorine
• Note: Hydrogen is a nonmetal even though it is located on
the periodic table where most of the metals are found.
© Stephanie Elkowitz 66 Atoms & Reactions
• Metalloids have characteristics in-between metals and
nonmetals. They are shiny and have a “metallic
appearance.” Although they look like metals, most
metalloids are brittle, not malleable and fair conductors.
• There are 6 commonly recognized metalloids: Boron,
Silicon, Germanium, Arsenic, Antimony and Tellurium.
• Astatine and/or Polonium are sometimes included. The
periodic table used in this presentation includes Polonium.
© Stephanie Elkowitz 67 Atoms & Reactions
• Noble gases are inert
(unreactive) substances.
• As their name suggests, Noble
gases are gases at room
temperature.
• Noble gases are poor
conductors of heat and
electricity.
• There are 6 Noble gases:
Helium, Neon, Argon, Krypton,
Xenon and Radon.
© Stephanie Elkowitz 68 Atoms & Reactions
• Recall: Elements in the same group have the same number
of valence electrons.
© Stephanie Elkowitz 69 Atoms & Reactions
• The number of valence electrons is easy to determine for
groups 1, 2 and 13 through 18.
© Stephanie Elkowitz 70 Atoms & Reactions
1 V
ALE
NC
E e
-
2 V
ALE
NC
E e
-
3 V
ALE
NC
E e
-
4 V
ALE
NC
E e
-
5 V
ALE
NC
E e
-
6 V
ALE
NC
E e
-
7 V
ALE
NC
E e
-
8 V
ALE
NC
E e
-
• There are exceptions to this pattern in the center block of
elements and to the rows of elements below the table. These
elements are called transitional metals.
• Transitional metals have complicated electron configuration.
Electrons are NOT always added to the outermost orbital.
© Stephanie Elkowitz 71 Atoms & Reactions
TRANSITIONAL METALS
TRANSITIONAL METALS
• Some groups have specific names because they share
similar properties:
– Group 1: Alkali Metals
– Group 2: Alkaline (Earth) Metals
– Group 3-12: Transitional Metals
– Group 17: Halogens
– Group 18: Noble Gases
• Elements in these groups share similar properties because
they have the same number of valence electrons.
© Stephanie Elkowitz 72 Atoms & Reactions
• The named groups are outlined in the period table below.
© Stephanie Elkowitz 73 Atoms & Reactions
HA
LO
GE
NS
NO
BLE
GA
SE
S
ALK
ALI M
ETA
LS
ALK
ALIN
E M
ETA
LS
TRANSITIONAL METALS
TRANSITIONAL METALS
• Alkali Metals are soft and shiny
metals found in Group 1.
• All Alkali Metals have one valence
electron. They readily lose this
electron to form a +1 ion.
• Hydrogen is NOT an Alkali Metal - it
is a nonmetal. It is placed in this
group because it has one valence
electron.
• Alkali Metals are extremely
reactive. If you put any of these
pure elements in water, they can
cause a huge explosion. © Stephanie Elkowitz 74 Atoms & Reactions
• Alkaline Earth Metals are Group 2
elements.
• These elements have two valence
electrons and readily lose these
electrons to form a +2 ion.
• Alkaline Earth Metals are the 2nd
most reactive elements.
• Within this group, you will find
Calcium. Calcium is important to
building your bones.
© Stephanie Elkowitz 75 Atoms & Reactions
• Transitional Metals are metallic elements found in Groups 3-
12 and in the rows below the periodic table as well.
• Transitional metals have complicated electron configurations.
They do not always completely fill a shell before beginning to
fill the next shell.
© Stephanie Elkowitz 76 Atoms & Reactions
• Halogens are nonmetals found in
Group 17.
• All of these elements have 7
valence electrons. They readily
gain one electron to form a -1 ion.
• Halogens are very reactive.
Fluorine is the most reactive
element in this group. The
reactivity of the elements in this
group decreases as you move
down the column.
© Stephanie Elkowitz 77 Atoms & Reactions
• Noble Gases are elements found
in Group 18.
• All of these elements are gases.
• Noble Gases have a full outer
shell. Except for Helium, all the
elements in this group have 8
valence electrons. Helium has a
full first shell with 2 valence
electrons.
• Noble Gases are very unreactive.
They rarely react with other
elements.
© Stephanie Elkowitz 78 Atoms & Reactions
• Below the periodic table are two rows of elements. These
rows make up the F Block of the periodic table.
• Elements in the F Block are part of the periodic table but it’s
easier to display them on the bottom of the table.
• All of the elements are transitional metals.
• Each row is considered a series of elements.
– The first row is called the lanthanide series.
– The second row is called the actinide series.
© Stephanie Elkowitz 79 Atoms & Reactions
• The Lanthanide Series includes chemical elements with
atomic numbers 57 through 71.
• All the elements are shiny, silvery metals.
• These elements are rare.
• Lanthanide elements are useful to superconductors, glass
production and lasers.
© Stephanie Elkowitz 80 Atoms & Reactions
• The Actinide Series includes chemical elements with atomic
numbers 89 through 103.
• All the elements are radioactive.
• Some of the actinide elements are synthetic - they are not
naturally found in Earth. Elements with atomic numbers 95
through 103 are synthetic elements.
• Plutonium is used in atomic weapons.
• Uranium has been used in atomic weapons. It is most often
used to produce electricity via nuclear power.
© Stephanie Elkowitz 81 Atoms & Reactions
• There are other synthetic elements in the periodic table.
• In total, there are 24 synthetic elements. These elements
have atomic numbers 95 through 118.
• Synthetic elements do not occur naturally on Earth.
• All synthetic elements were made in a laboratory.
• All synthetic elements are unstable and radioactive. They
decay rapidly, some in only a few hundred microseconds.
• Some elements, such as Technetium and Plutonium, are
synthetically made. However, they are not purely synthetic.
They exist naturally in very small quantities on Earth.
© Stephanie Elkowitz 82 Atoms & Reactions
• There are patterns to the properties of elements in the periodic table.
These patterns are called trends.
© Stephanie Elkowitz 83 Atoms & Reactions
• Atomic radius is a measure of the size of an
atom.
• Specifically, atomic radius is a measure of
how “wide” an atom is from the center of its
nucleus to the outermost electron shell.
© Stephanie Elkowitz 84 Atoms & Reactions
Atomic radius
• As you move down a column, atoms have more orbitals.
Orbitals increase the size of an atom. Therefore, atomic
radius increases as you move down a column.
• As you move across a row, atomic radius decreases. As an
atom gains electrons in the same orbital, the electrons are
more attracted to and “pulled” towards the positive nucleus.
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• Atoms with nearly a full valence electron shell want to gain
electrons to completely fill the shell. These atoms have a
high electronegativity. Electronegativity is the tendency of an
atom to attract an electron. The greater the electronegativity,
of an atom, the mire likely the atom will attract an electron.
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• As you move across a row, electronegativity increases. This
makes sense because atoms have more valence electrons.
• As you move down a column, an element’s electronegativity
decreases. It decreases because larger atoms have a more
difficult time attracting electrons.
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• Ionization energy is the energy required to remove an electron
from an atom. The greater the ionization energy of an atom,
the “harder” it is to remove an electron from the atom.
• Atoms with 1 or 2 valence electrons easily give up electrons
while atoms with a nearly a full valence electron shell do not.
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• As you move across a row, ionization energy increases. Atoms
have more valence electrons and do not want to give them up.
• As you move down a column, ionization energy decreases.
The size of an atom increases so it’s easier for an atom to
lose electrons.
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• Elements, except hydrogen, on the left side of the period table
are metals. These elements have metallic character.
• Metallic character refers to the characteristics of metals, such
as luster, malleability, ductility and conductivity.
• Metalloids have some metallic character. Nonmetals do not
have metallic character, which is why they are called nonmetal.
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• As you move across the periodic table (left to right),
elements lose metallic character. In other words, metallic
character decreases.
• Elements transition from metals to metalloids to nonmetals
as you move across the table. Metals have the most
metallic character. Nonmetals have no metallic character.
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